Question

EQUILIBRIA PROBLEM

Answer 1

\[AgBr_{(s)}<--> Ag^+_{(aq)} + Br^-_{(aq)} ~~~~~ K_{sp}=5 \times 10^{-13}\]
i) Calculate [Ag+] in a saturated aqueous solution of AgBr.
ii) Silver ions form complexes with ammonia and with amines.
\[Ag^+_{(aq)}+2NH_3(aq)\leftarrow \rightarrow [Ag(NH_3)_2]^+(aq)\]
Using expression for Kc calculate the [NH3] needed to change the [Ag+] in a 0.10 moldm-3 solution of silver nitrate to the value that you calculated in (i)

Answer 2

the first part i have calculated is
(i) Ksp=[Ag+]^2
\[[Ag^+]= \sqrt{Ksp}=7.07\times10^{-7}\]

Answer 3

For the second part, they given the Kc is \(\Large K_c=1.7 \times 10^{7}\)

Answer 4

my question is the second part, where does silver nitrate comes into play?

Answer 5

the silver nitrate is just the source of silver ions. the nitrate ions don't do anything in this equlibrium, we can effectively ignore them.

Answer 6

so which value of [Ag+] should i use? should i use 0.1moldm-3 or 7.07x10-7moldm-3

Answer 7

you actually need both. This is a stoichiometry problem at the moment, not necessarily an equilibrium problem. You need to know how muc hNH3 to add in order to CHANGE the concentration of Ag+ ions from 0.1M to 0.1x10-7M

Answer 8

should i create an equation then solve the stoichiometry?

Answer 9

\[AgNO_3 \rightarrow Ag^+ + NO_3^-\]

Answer 10

@Calculator bro it's been ages since i left ionic eullibria but i must say you are going in a right direction, use stoichiometry and then use Ksp for that

Answer 11

how do you work out the stoichiometry? they are both concentrations

Answer 12

\[K _{sp}=\frac{ [Ag+][Br-] }{ [AgBr] }\] just take the concentrations of the ions

Answer 13

hmm, maybe i need an ICE table

Answer 14

yea i think so

Answer 15

an ICE table (or RICE box, when I teach it) would be perfect here.